Problem 57 - Chemistry and Chemical Reactivity

Question

In the following reactions, decide which reactant is oxidized and which is reduced. Designate the oxidizing agent and the reducing agent. (a) \(\mathrm{C}_{2} \mathrm{H}_{4}(\mathrm{g})+3 \mathrm{O}_{2}(\mathrm{g}) \rightarrow 2 \mathrm{CO}_{2}(\mathrm{g})+2 \mathrm{H}_{2} \mathrm{O}(\ell)\) (b) \(\mathrm{Si}(\mathrm{s})+2 \mathrm{Cl}_{2}(\mathrm{g}) \rightarrow \mathrm{SiCl}_{4}(\ell)\)

Step-by-Step Solution

Answer

(a) Oxidized: \( \mathrm{C}_{2} \mathrm{H}_{4} \), Reducing agent: \( \mathrm{C}_{2} \mathrm{H}_{4} \); Reduced: \( \mathrm{O}_{2} \), Oxidizing agent: \( \mathrm{O}_{2} \). (b) Oxidized: \( \mathrm{Si} \), Reducing agent: \( \mathrm{Si} \); Reduced: \( \mathrm{Cl}_{2} \), Oxidizing agent: \( \mathrm{Cl}_{2} \).

Step 1: Identify Oxidation and Reduction in Reaction (a)

In the reaction \( \mathrm{C}_{2} \mathrm{H}_{4}(\mathrm{g})+3 \mathrm{O}_{2}(\mathrm{g}) \rightarrow 2 \mathrm{CO}_{2}(\mathrm{g})+2 \mathrm{H}_{2} \mathrm{O}(\ell) \), observe the change in oxidation states. The carbon in \( \mathrm{C}_{2} \mathrm{H}_{4} \) goes from an oxidation state of -2 to +4 in \( \mathrm{CO}_{2} \), indicating oxidation. Oxygen goes from 0 in \( \mathrm{O}_{2} \) to -2 in \( \mathrm{CO}_{2} \) and \( \mathrm{H}_{2} \mathrm{O} \), indicating reduction.

Step 2: Identify Agents in Reaction (a)

The substance that gets oxidized, \( \mathrm{C}_{2} \mathrm{H}_{4} \), acts as the reducing agent because it donates electrons. Conversely, \( \mathrm{O}_{2} \), which is reduced, acts as the oxidizing agent because it accepts electrons.

Step 3: Identify Oxidation and Reduction in Reaction (b)

In the reaction \( \mathrm{Si}(\mathrm{s})+2 \mathrm{Cl}_{2}(\mathrm{g}) \rightarrow \mathrm{SiCl}_{4}(\ell) \), silicon goes from an oxidation state of 0 to +4, indicating oxidation. Chlorine goes from 0 in \( \mathrm{Cl}_{2} \) to -1 in \( \mathrm{SiCl}_{4} \), indicating reduction.

Step 4: Identify Agents in Reaction (b)

In this reaction, \( \mathrm{Si} \), which is oxidized, is the reducing agent. \( \mathrm{Cl}_{2} \), which is reduced, acts as the oxidizing agent.

Key Concepts

Oxidizing Agent

In every redox reaction, there is a substance that gets reduced. This substance is known as the oxidizing agent. It might sound a bit counterintuitive because the oxidizing agent is not the one being oxidized, but it is crucial to remember that it facilitates the oxidation of another substance. Here's how it works:

When a chemical species acts as an oxidizing agent, it gains electrons from the substance being oxidized. By accepting electrons, it undergoes reduction. This is why, despite its name, the oxidizing agent itself becomes reduced.

If we consider the reaction \[ \mathrm{C}_{2} \mathrm{H}_{4}(\mathrm{g})+3 \mathrm{O}_{2}(\mathrm{g}) \rightarrow 2 \mathrm{CO}_{2}(\mathrm{g})+2 \mathrm{H}_{2} \mathrm{O}(\ell) \]we see that oxygen (\(\mathrm{O}_{2}\)) is the oxidizing agent. It accepts electrons from \(\mathrm{C}_{2} \mathrm{H}_{4}\) and, as a result, gets reduced to form \(\mathrm{H}_{2} \mathrm{O}\) and \(\mathrm{CO}_{2}\).

In another example of the reaction \( \mathrm{Si}(\mathrm{s})+2 \mathrm{Cl}_{2}(\mathrm{g}) \rightarrow \mathrm{SiCl}_{4}(\ell) \), we find that chlorine \(\mathrm{Cl}_{2}\) is the oxidizing agent. By gaining electrons from silicon (\(\mathrm{Si}\)), it undergoes reduction. Remember, oxidizing agents are essential players in the process of electron transfer.

• Facilitates oxidation
• Undergoes reduction
• Accepts electrons

Reducing Agent

Just as there is an oxidizing agent, there is also a reducing agent in redox reactions. The reducing agent is the substance that gets oxidized. This means it donates electrons to another substance and itself remains oxidized.

The name might suggest it reduces something, and indeed it does: it reduces the other reactant by supplying the electrons that allow the reduction to occur.

In the reaction \( \mathrm{C}_{2} \mathrm{H}_{4}(\mathrm{g})+3 \mathrm{O}_{2}(\mathrm{g}) \rightarrow 2 \mathrm{CO}_{2}(\mathrm{g})+2 \mathrm{H}_{2} \mathrm{O}(\ell) \), \(\mathrm{C}_{2} \mathrm{H}_{4}\) takes on the role of the reducing agent. It donates electrons to \(\mathrm{O}_{2}\), allowing \(\mathrm{O}_{2}\) to be reduced while \(\mathrm{C}_{2} \mathrm{H}_{4}\) itself is oxidized to \(\mathrm{CO}_{2}\).

Similarly, in the reaction \( \mathrm{Si}(\mathrm{s})+2 \mathrm{Cl}_{2}(\mathrm{g}) \rightarrow \mathrm{SiCl}_{4}(\ell) \), silicon \(\mathrm{Si}\) acts as the reducing agent. It donates electrons to chlorine, permitting it to be reduced to chloride ion in \(\mathrm{SiCl}_{4}\).

• Facilitates reduction
• Undergoes oxidation
• Donates electrons

Oxidation States

Oxidation states, sometimes referred to as oxidation numbers, are a way to keep track of how many electrons an atom has gained, lost, or shared during a chemical reaction. Understanding oxidation states is crucial for identifying which elements are oxidized and reduced in a reaction.

**Determining Oxidation States:**

• For any pure element, the oxidation state is zero. This includes diatomic molecules like \(\mathrm{O}_{2}\) and \(\mathrm{Cl}_{2}\).
• For a simple ion, its oxidation number equals the charge of the ion.
• In compounds, hydrogen usually has an oxidation state of +1, and oxygen is typically -2.

Consider the reaction \( \mathrm{C}_{2} \mathrm{H}_{4}(\mathrm{g})+3 \mathrm{O}_{2}(\mathrm{g}) \rightarrow 2 \mathrm{CO}_{2}(\mathrm{g})+2 \mathrm{H}_{2} \mathrm{O}(\ell) \). Here:

• Carbon in \(\mathrm{C}_{2} \mathrm{H}_{4}\) starts with an oxidation state of -2, changing to +4 in \(\mathrm{CO}_{2}\). This change signifies oxidation.
• Oxygen in \(\mathrm{O}_{2}\) goes from an oxidation state of 0 to -2 in \(\mathrm{CO}_{2}\) and \(\mathrm{H}_{2} \mathrm{O}\), indicating reduction.

For another instance, take the reaction \( \mathrm{Si}(\mathrm{s})+2 \mathrm{Cl}_{2}(\mathrm{g}) \rightarrow \mathrm{SiCl}_{4}(\ell) \):

• Silicon goes from an oxidation state of 0 to +4, showing its oxidation.
• Chlorine switches from 0 in \(\mathrm{Cl}_{2}\) to -1 in \(\mathrm{SiCl}_{4}\) as it is reduced.

Being able to determine oxidation states lets you track Electron transfer and predict the roles of oxidizing and reducing agents effectively.